Rajasthan Board RBSE Class 10 Science Notes Chapter 7 Atomic Theory, Periodic Classification, and Properties of Elements
Dalton’s Atomic Theory:
- Each substance is made up of small particles which are called atoms.
- Atoms are indivisible.
- All atoms of an element have same properties, e.g. mass, size, chemical property, etc.
- Atoms of different elements have different properties, e.g. mass, size, chemical property, etc. -i Atoms of different elements combine in whole number ratios to form molecules.
- Atoms get rearranged during chemical reactions. A chemical reaction can neither create an atom nor destroy an atom.
Thomson’s Atomic Model:
The atomic model of Thomson is known as plum pudding model. According to this, atom is very small in size; 10-10m. An atom is like a pudding which is positively charged. Negatively charged electrons are embedded in this pudding, the way plums are embedded in pudding. Thomson’s model was successful in explaining electrical neutrality of an atom.
Rutherford’s Gold Foil Experiment:
Rutherford took a very thin gold foil with 100 nm thickness. He bombarded the gold foil with alpha rays. It is important to note that alpha particles are positively charged.
Observations made by Rutherford:
- Almost all the alpha rays passed through the gold foil without any deviation in path.
- Some of the alpha rays showed slight deviation
- A very few alpha radiations bounced back to their source.
Rutherford’s Conclusions:
- A major portion of atom is hollow and bears no charge. Hence, most of the alpha rays passed through without deviation.
- Some of the alpha-rays showed slight deviation means that there is a positively charged portion in atom. The positive charge in atom offered repulsive force on alpha particles.
- Very few alpha rays bounced back means that the positively charged portion’s size is very small as compared to the overall size of atom. Diameter of positive charged portion i.e., nucleus is around 10-15m.
Rutherford’s Atomic Model:
- The whole positive charge and mass of an atom lie in its nucleus.
- Most of the atom is hollow in which electrons move around the nucleus. The circular path on which electrons move is called orbit.
- Since an atom is electrically neutral hence number of electrons in an atom is same as the number of protons.
Drawbacks of Rutherford’s Model:
- This model was unable to explain the stability of an atom.
- This model could not explain electronic configuration of an atom.
- According to Maxwell’s theory, an electron moving in a circular orbit would experience a loss in
energy. As a result, the electron would begin to move on a spiral path and would finally fall on
the nucleus. This should result in collapse of the atom. But no atom collapses in reality.
Hypothesis of Neil’s Bohr
- The nucleus of an atom is positively charged and protons are present in the nucleus.
- Electrons move around the nucleus in orbits of finite energy and radius. These orbits are also called energy levels.
- These orbits are concentric and are denoted by n. The value of n is expressed in whole numbers,
e.g. 1, 2, 3,4, 5, etc. They are also denoted by letters K, L, M, N, - A higher number of n means the orbit is farther from the nucleus and has higher level of energy. The first orbit or K or n = 1 has lowest level of energy.
- Angular momentum of electron in an orbit is mvr = h/2Π where h = Plank’s constant, m = mass of electron, v = velocity of electron and r = radius of orbit. Hence, electron can only move in an orbit in which the angular momentum is ng/2n.
- When an electron continues to move in a definite orbit, there is no loss in its energy.
- When an atom gets energy from any external source then an electron jumps to a higher energy level. When an electron loses some energy, it jumps down to a lower energy level.
Drawbacks of Bohr’s Model:
- This model fails to explain an atom with a higher number of electrons.
- This model does not explain formation of chemical bond.
- High resolution spectrum of atom shows more than one spectral line which cannot be explained by this model.
Dobereiner’s Classification:
Dobereiner tried to classify and group elements in the order of increasing atomic masses, he made group of three elements (triads) having similar chemical properties. The atomic mass of the middle element of the triad being equal to the arithmetic mean of the atomic masses of the other two elements.
Triads: Li, Na, K
6.9,22.9,39
\(6.9+39=\cfrac { 45.9 }{ 2 } =22.95\quad \)
Drawbacks of Dobereiner’s Triads
- He could make only three triads. ‘
- At that time only few elements (54) were known, the grouping was not done for all the elements so, it was not accepted.
- Newland’s Law of Octaves: He grouped seven elements and arranged them in increasing order of their atomic masses. He started with lowest element hydrogen and ended with thorium. Every eighth element showed the properties similar to that of first. It was called the Law of Octaves, (like music notes the first and the eighth note is same, similar to octaves.)
Drawbacks of Newland’s Octaves
- He could group the elements only till calcium, after Ca the remaining elements did not follow the law of Octaves.
- Newland assumed that only 56 elements existed in nature and no new element will be discovered.
- He tried to jumble few elements to make possible that the Law of Octaves is followed, i.e., the properties of 1st and 8th element is same.
- In order to fit the elements in his octave, he adjusted two elements, cobalt and nickel by placing in the same group of fluorine, chlorine and bromine. Iron which resembled cobalt and nickel was placed in some other group.
Mendeleev’s Periodic Law:
The properties of the elements are the periodic functions of then atomic masses.
- Mendeleev arranged the elements in the increasing order of their masses and the elements lying in the group have same properties.
- 63 elements were known at that time.
- He made a table of 8 groups (vertical columns) and 6 periods (horizontal rows).
Achievements:
While arranging the elements Mendeleev left 3 gaps and he predicted 3 elements as Eka boron, Eka aluminium and Eka silicon. He also predicted the properties of these elements.
All the 3 elements were discovered later; they were named as Scandium, Gallium and Germanium.
Noble gases were not discovered then. After discovery they got a separate group without disturbing
the elements in the table.
Limitations/Anomalies of Mendeleev’s Classification
- While arranging the elements in this table, he placed cobalt (at. mass 58.9) before nickel (at. mass 58.7).
- The position of hydrogen was not clear because it was placed along with the alkali metals.
- It could have been placed along with halogen family because it forms di-atomic molecule and like them it combines with metals forming covalent bonds.
- Isotopes have different atomic mass but they were not given any separate position in this table.
Modern Periodic Law:
Properties of the elements are periodic functions of their atomic number.
- Modern Periodic Table: It has 7 horizontal rows called periods and 18 vertical columns called groups.
- Period number of an element = Number of shells.
- It is divided in four blocks viz. S-block (alkali and alkaline earth metals), p-block (main group elements), d-block (transition elements) and f-block (inner transition elements).
Periodic Properties
Atomic size: Atomic size is the distance between the outermost orbit and the centre of the nucleus. On moving down the group, the atomic size of elements goes on increasing, as one shell get added as we move from top to bottom. On moving from left to right the atomic size decreases. This is due to increase in nuclear charge that pulls the electrons towards the nucleus.
- Ionic radius: When an atom donates its electron it forms cation which due to disappearance of outer shell or increase in effective nuclear charge has less radius than neutral atom. While when an atom accepts an electron and forms anion which has larger radius than neutral atom due to opposite effect as that of cations.
- van der waal’s radius is greater than covalent radius.
Ionisation enthalpy:
It increases on moving from left to right across a period due to increase in effective nuclear charge. It decreases on moving top to bottom along a group due to increase in atomic size and decrease in effective nuclear charge.
Electron gain enthalpy:
On moving from left to right atomic size decreases and effective nuclear charge increases thus, electron gain enthalpy becomes more negative. On moving down the group, atomic size increases thus, electron gain enthalpy becomes less negative.
Electronegativity:
On moving from left to right atomic size decreases thus, electronegativity increases while it decreases on moving from top to bottom in a group.
Valency:
Valency is constant for all elements in a group. Along a period from left to right, valency increases from in to 4 and then reduces to zero.
Elements of d-block and f-block show variable valency.
Metallic character:
On moving down the group, the metallic character goes on increasing due to increase in atomic size. On moving from left to right, across a period the metallic character decreases and the non-metallic character increases due to decrease in atomic size.
Chemical Reactivity:
The chemical reactivity of metals increases down the group as the metals readily lose electron as the atomic size increases down the group.
In case of non-metals, the reactivity decreases as we move down the group, the atom at the top of the group readily accepts electron.